Holding molecules together - van der Waals forces

A while back, I did a post for Chemistry week about hydrogen bonds. In it, I mentioned why I find intramolecular forces so fascinating; they are interactions on such a tiny scale that hold together everything from small molecules like water to massive molecules like the enzymes and multi-enzyme complexes that I study.

The hydrogen-bond post seems to pick up a fair amount of visitors each month, and I like to think it has been a useful post for those studying Chemistry for school, or trying to remember their schoolwork at university level. In the spirit of this I thought I would write another post on intramolecular forces, covering one of my favourite forces - van der Waals forces of attraction. They are also known as 'induced dipoles' but van der Waals is more fun to say.

Hydrogen bonds (covered here) are often thought of as fairly weak, but they are solid iron-and-concrete compared to van der Waals forces. Like hydrogen bonds van der Waals rely on dipoles, a difference in charge between two molecules. But unlike hydrogen bonds the van der Waals dipole is not permanent, but transient.

A dipole relies on one side of the molecule having more electrons around it (and therefore being slightly negative) while the other side has consequently fewer electrons (and is therefore positive). That's fine for molecules like water which have a dipole, but for molecules like, say, iodine which is composed of two iodine atoms sharing electrons, it's a problem. In iodine all the electrons are shared equally between the two atomic centres and there is no dipole. But iodine can form a solid at room temperature, which requires pretty strong forces between molecules to achieve.

To be fair it's not a particularly stable solid ... you can see the iodine crystals here starting to turn into a purple vapour (sublimation). Image from wikimedia.

Iodine is a big atom with lots of electrons (53 to be exact). This means that in an iodine molecule there will be 106 electrons all whizzing about at high speed, many of which are far away enough from the central nucleus to be a lot less rigid in exactly where they are around the molecule. Which means that occasionally, just for a brief moment, there will be more electrons on one side of the molecule than the other ...

Instant dipole!

An illustration of the induced dipole around an iodine molecule. Each little pink line represents one electron. I have exaggarated the disproportionate number of electrons around each nucleus in order to illustrate the concept - in 'reality' the effect would be less obvious. Image (c) me.

When a dipole forms on one molecule it will start effecting molecules around it. A build-up of electrons on one side of a molecule forms a slight negative force, which will repel electrons on nearby molecules, making them slightly positive on one side and thus propagating the dipole. The negative side of one molecule can then form weak bonds with the positive side of the neighbouring one.

This doesn't last, because the elecrons are all still whizzing around at high speed and it isn't wonderfully energetically favourable for them to be mostly on one side of the molecule. It's a transient dipole, and a transient force, but when you have millions of molecules making millions of connections enough force is generated to hold the iodine molecules in a (mostly) solid state.

These van der Waals forces don't work for every molecule though. Iodine has a large number of electrons that are an appreciable distance from the nucleus but other molecules, such as fluorine, do not. Fluorine has only 9 electrons per atom and these are held tightly in place by the central nucleus. They'll still be whizzing around very fast, but it will be along far more rigidly constrained pathways. Fluorine therefore will not feel the effects of van der Waals forces.

And that is why fluorine is a gas! Image from wikimedia commons.

The views expressed are those of the author and are not necessarily those of Scientific American.

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